Atomic Structure and Bonding

Atomic Structure and Bonding

• An atom, the smallest unit of an element, retains the properties of that element and consists of a nucleus surrounded by electrons
• Protons are positively charged particles in the nucleus that determine an element's atomic number
• Neutrons are neutral particles in the nucleus that contribute to the mass of an atom along with protons
• Electrons are negatively charged particles orbiting the nucleus and are involved in chemical bonding and reactions
• The nucleus is the dense, positively charged central part of an atom containing protons and neutrons
• Atomic number is the number of protons in the nucleus, defining the element's identity and position on the periodic table
• Mass number is the total count of protons and neutrons in an atom's nucleus, indicating its mass
• Valence electrons are located in the outermost shell and determine how an atom will react chemically
• Electron configuration is the arrangement of electrons in an atom's shells, which helps predict an element's chemical properties
• The periodic table organizes known elements by increasing atomic number and groups them based on similar chemical properties
• A group is a vertical column, where elements share chemical properties and the same number of valence electrons
• A period is a horizontal row, where elements have the same number of electron shells
• Metals are shiny elements, conduct heat and electricity well, and are malleable and ductile; found on the left side of the periodic table
• Non-metals are typically not shiny, poor conductors of heat and electricity, and are brittle in solid form; found on the right side of the periodic table
• Metalloids possess properties of both metals and nonmetals and are located along the zig-zag line on the periodic table
• Electron configurations can be found using the 2,8,8 method, filling electron shells with 2, 8, and 8 electrons, then subtracting 2 from the third shell if needed to total to 18

Atomic Stability

• Atoms form bonds to achieve a stable electron configuration, often resembling that of noble gases
• Atoms gain stability by sharing, gaining, or losing electrons via chemical bonds to fill their outer electron shells
• This process lowers the potential energy and increases stability
• In a water molecule, oxygen and hydrogen share electrons to complete their outer shells, resulting in a stable compound

Organization of the Periodic Table

• Groups are the vertical columns in the periodic table and elements in the same group exhibit similar chemical properties due to having the same number of valence electrons
• Periods are the horizontal rows, and moving across a period sees increasing atomic numbers with a transition from metallic to nonmetallic properties
• Metals reside on the left side and in the middle, generally shiny, conductive, and malleable
• Nonmetals are found on the right side, often not shiny, poor conductors, and brittle
• Metalloids have properties of both metals and nonmetals, located along the zigzag line dividing the table
• The modern periodic table arranges elements by increasing atomic number rather than atomic mass
• Electron configuration describes the distribution of electrons in an atom's orbitals, influencing similar chemical properties within a group

Metals

• Metalloids have intermediate properties between metals and non-metals having better electrical conductivity than non-metals, but less than metals' electrical conductivity, like silicon and arsenic
• A bond is a force holding atoms together in a molecule or compound, including ionic, covalent, and metallic types
• A cation is a positively charged ion formed when an atom loses electrons; sodium becoming a sodium cation (Na+) after losing an electron is an example
• Delocalized electrons are not associated with a single atom or covalent bond; these electrons move freely in metals, enabling electrical conductivity
• Electrostatic attraction is the force drawing oppositely charged particles together, such as cations and anions in ionic bonds
• A lattice is a regular arrangement of atoms, ions, or molecules in a solid structure
• Malleability is a material's ability to be hammered or rolled into thin sheets without breaking, a property common in metals
• Ductility means to be stretched into a wire without breaking, a characteristic shown by metals
• Electrical conduction is the capacity to allow electric current flow that is high in metals due to delocalized electrons
• Thermal conductivity is how well a material conducts heat, materials with high thermal conductivity transfer heat rapidly

Structure and Properties of Metals

• Metal atoms arrange in a crystal lattice structure like a 3D grid with atoms closely packed together
• Bonding involves electrons moving freely to form a "sea of electrons" that enables electrical conductivity
• Common lattice structures include body-centered cubic (BCC), face-centered cubic (FCC), and hexagonal close-packed (HCP) which gives metals their unique properties
• Malleability and ductility result from metal atoms being able to slide past each other allowing the metal to be shaped
• Conductivity occurs via free-moving electrons that enable efficient electricity and heat conduction
• Luster comes from free electrons reflecting light, resulting in a shiny appearance
• Strength is from strong bonds creating durable metal structures

Reactivity of Metals

• Group 1 alkali metals (Lithium, Sodium, Potassium) are very reactive especially with water
• Alkali metals have one electron in their outermost shell (e.g., Li: 2,1; Na: 2,8,1)
• Reactivity in Group 1 increases down the group due to the outer electron being more easily lost as it's further from the nucleus
• Group 2 alkaline earth metals (Beryllium, Magnesium, Calcium) are reactive, though less so than Group 1 metals
• Alkaline earth metals have two electrons in their outermost shell (e.g., Be: 2,2; Mg: 2,8,2)
• Reactivity in Group 2 increases too as you move down, but less reactive than Group 1 metals
• Reactivity increases down each group due to increasing atomic size and shielding, making outer electrons easier to shed as they grow weaker
• Outer electrons dictate reactivity; Group 1 metals readily lose one electron, while Group 2 metals lose two

Ionic Bonds

• An ion is an atom or molecule with a net electric charge resulting from loss or gain of electrons
• An anion is a negatively charged ion from having gained electrons
• A polyatomic ion is a charged entity made of covalently bonded atoms
• An ionic bond involves electrostatic attraction between oppositely charged ions, often between a metal and a nonmetal
• Ionic bonds occur as metals lose electrons to become cations and nonmetals gain electrons to become anions
• Cations and anions attract each other and form a stable ionic bond
• Ionic compounds organize in a 3-D crystal lattice by maximizing attraction and minimizing repulsion
• Lattice structure contributes to the stability of ionic compounds, reducing likelihood of breakdown under normal conditions
• Ionic compounds have high melting and boiling points due to strong electrostatic attraction between ions
• They are hard but brittle because the rigid lattice structure may cause ions to repel if displaced
• Electrical conductivity happens only when ionic compounds are melted or dissolved in water due to ions being free and not located in the lattice
• Many ionic compounds are soluble in water because it separates ions from the lattice

Ionic Compounds

• Calcium fluoride (CaF2) is formed by combining calcium (+2 charge) and fluoride (-1 charge)
• To formulate ionic compounds, balance total positive and negative charges
• Sodium chloride (NaCl) is formed by sodium (+1) and chloride (-1)
• Magnesium oxide (MgO) is formed by magnesium (+2) and oxide (-2) ions

Covalent Bonds

• Covalent bonds form when atoms share electrons, promoting stability
• A covalent bond involves atoms sharing electrons, typically between nonmetals
• Lone pairs are electrons not involved in bonding, residing on a single atom and influencing its shape and reactivity
• A covalent molecule is formed by covalent bonds and oxygen (O2)
• A covalent network is a large configuration of covalently bonded atoms, such as diamond, with unique physical properties
• Nonmetals typically have 4-8 valence electrons that affect bonding
• Nonmetals gain or share electrons to fill outer shell, achieving stability
• Bond behavior includes forming ionic bonds with metals and covalent bonds with other nonmetals

Covalent and Nonmetals

• Nonmetals form covalent bonds for a stable configuration located at the right side of the periodic table (excluding noble gasses)
• Nonmetals share electrons to gain stability rather than transferring due to lacking metals to take electrons from
• Bond formation happens when two nonmetals share one (single), two (double), or three (triple) pairs of electrons
• Types of covalent bonds include single, double, and triple bonds like hydrogen(H2, single), oxygen (O2, double), and nitrogen (N2, triple)

Properties of Covalent Molecules

• Covalent molecules have definite physical properties because of bonding and structure
• Sharing electrons leads to shapes like linear or tetrahedral
• Molecules can be polar (uneven charge distribution) or nonpolar (even distribution), affecting solubility
• Lower than ionic compounds are melting/boiling points due to weaker intermolecular forces
• Polar molecules dissolve in polar solvents, while nonpolar dissolve in nonpolar
• Covalent compounds do not have electrical conductivity because the lack free-moving charged particles

Covalent Structure

• Ammonia(NH3), methane (CH4), water (H2O), oxygen (O2), nitrogen (N2), chlorine (Cl2), bromine (Br2), iodine (I2), and fluorine (F2) are examples of covalent molecules

Diamond Properties

• Diamond's structure is a tetrahedral structure where each carbon atom covalently bonds to four others, making a 3D network
• Its strong covalent bonds make it the hardest natural substance
• Transparency and a high refractive index make Diamonds transparent
• Diamond is an excellent electrical insulator and a good thermal conductor
• Structure defines the characteristics of graphite with carbon atoms bonded to three others in hexagonal rings

Graphite

• Graphite layers are held together by weak van der Waals forces, allowing layers to easily slide across each other making it soft and slippery
• Electrical conductivity is present in the presence of free electrons within the layers
• Graphite is opaque with a metallic luster
• Diamond has strong bonds and is hard to break
• Graphite has a strong but weak connection

Electron Configuration

• An atom contains a positively charged nucleus with neutrons and protons for its mass, with electrons orbiting in shells
• Electrons follow a 2,8,8 arrangement which designates charge
• Valence electrons in the outer shell delocalized electrons move between atoms
• Charge 1 mass 0 electron charge -1 mass -1